The present invention relates in general to aspirin, and in particular to a new and useful composition and method for rapidly soluble aspirin.
The important role played by aspirin in the treatment and management of the two most common major diseases afflicting the world's population, i.e. arthritis and heart attacks, is well known. It is an effective analgesic and anti-inflammatory agent. In addition, its beneficial effect on the immune system makes aspirin a promising drug for the treatment of cancer, AIDS, cataracts, allergies and other diseases in which the immune system is involved. These valuable pharmacological properties makes aspirin the most widely used drug in the world.
Aspirin possesses side-effects, however, the true severity of which has been recognized only recently. Being a sparingly soluble substance (0.33 gm in 100 ml of water) particles of aspirin adhere to gastrointestinal mucosa causing lesions, gastric and duodenal ulcers and often massive bleeding and death. These side effects are persistent and cumulative and occur in nearly all patients using aspirin therapy. Gastroscopic and clinical studies confirm the topical nature of these side effects. (Lancet 2:1222, 1938; 1:539, 1959; August 1980; Brit. Med. J. 2:7, 1955; New England J. Medicine 258:219, 1958; The Amer. J. of Digestive Diseases, 1961; Pharmacology, 25, 1982; The Annals of Internal Medicine, Sept. 1988).
Corrosive effects of aspirin on gastrointestinal mucosa were recognized early in the history of aspirin therapy and attempts to produce soluble forms of this drug were made continuously since then. See U.S. Pat. No. 740,703 of 1903. These efforts were concentrated almost exclusively on the preparation of various soluble salts: lithium, sodium, potassium, calcium, magnesium and with organic amines and amino acids (lysine, ornithine). The main disadvantages of these salts was that, in contrast to aspirin itself, they were not stable. Most of them contain water of crystallization which produces an intramolecular hydrolysis resulting in the splitting of the molecule into salicylic and acetic acids. Attempts to prevent this decomposition by removing the water and forming anhydrous salts result in hygroscopic products which are difficult to handle, expensive to produce and require special expensive moisture-proof packaging of each individual dose. The cost of such products was very high and they could not compete with aspirin commercially. In addition, the use of these salts often involves the ingestion of undesirably high amounts of metallic elements.
More recently, attempts were made to obviate these disadvantages and side effects by simpler and less expensive means. Thus, aspirin tablets were coated with layers of buffering agents designed to neutralize the gastric acidity. Calcium carbonate, magnesium carbonate, magnesium hydroxide, sodium aluminum carbonate, sodium aluminum glycinate and the like, are thus in use in a number of commercial aspirin products. Clinical studies show that these methods are not effective, partly because it is not possible to coat a tablet with sufficient amounts of buffering agent to neutralize the gastric acid. Even if it were possible to do so, there is an immediate natural response which causes the production of more acid, often in larger amounts than originally present (acid rebound). But, more importantly, buffering agents do not prevent the insoluble particles of aspirin from adhering to the gastro-intestinal mucosa and causing corrosion. Another method used in commercial tablets is to entero-coat them to prevent the aspirin release in the stomach and to exert its effect in the intestine. This simply results in shifting the locus of side-effects from the stomach to the intestine.
As mentioned above, the manufacture of soluble salts of aspirin is complicated and costly. A typical example is the manufacture of the sodium salt of aspirin as disclosed in U.S. Pat. No. 3,985,792. The first step of this process consists in reacting aspirin with sodium bicarbonate in water, which forms a solution of the sodium salt. In the second step, this solution is treated with isopropanol and cooled to 5.degree. C., which causes the crystallization of sodium acetylsalicylate dihydrate. The third step involves the filtration and the washing to the dihydrate. In the fourth step, the dihydrate, which is unstable, is dehydrated as soon as possible in a vacuum dryer or in a current of dry, inert gas, such as nitrogen. The final product is hygroscopic and must be handled, stored and packaged in humidity-controlled rooms. And finally, isopropanol, which is used in large amounts (10 lbs per lb of product) must be separated from water and recovered by fractional distillation. This involved process is further complicated by the fact that aspirin is decomposed by water and by isopropanol, which affects the yield, the purity and the stability of the final product.
Aspirin readily reacts with sodium bicarbonate to form an aqueous solution of the sodium salt. It would therefore appear that the costly and complicated preparation of the sodium salt could be avoided by simply adding aspirin and sodium bicarbonate, in pre-measured amounts corresponding to the desired dose, to a glass of water, stirring the mixture until the aspirin is dissolved and drinking the solution. This would provide aspirin solutions simply and inexpensively, which are free of corrosion-causing insoluble particles to patients on aspirin therapy. And, indeed, such products are commercially available, usually in the form of effervescent tablets containing a mixture of aspirin, sodium bicarbonate and citric acid.
The major drawback of such products is that in order to accomplish dissolution even of the smallest adult dose (325 mg, 5 grains) it was found necessary to use large amount of sodium bicarbonate (1900 mg), which represents a very large excess, since the theroretical amounts needed is only 152 mg. Even allowing for the fact that some of the sodium bicarbonate is neutralized by the citric acid, the amount of the bicarbonate present is equivalent to nearly 40 moles when only one mole is required for the reaction.
Aspirin is usually taken in dosages to two tablets 325 mg each, three times a day. The use of commercial tablets described above would, while giving particle-free solutions of aspirin, involve ingesting 3,000 mg of elemental sodium per day. This amount of sodium is considered medically detrimental to health in general, but particularly to older patients, and patients with hypertension. As a result, such products are used only for the relief of occasional minor pain or an upset stomach. They are never used on a regular basis by patients with arthritic or those on restricted sodium diets.
In the laboratory or during industrial manufacture, the amount of sodium bicarbonate used is 46.7 parts per 100 parts of aspirin, whereas in soluble aspirin tablet described above, the amount is 1250 parts, or about 20 times larger. The reason why it is necessary to use such large amounts of sodium bicarbonate in commercial tablets can be explained by the kinetics of the reaction involved.
In the preparation of sodium aspirin, whether in the laboratory or on an industrial scale, the amount of water used is as small as practically possible. Thus, for 100 parts of aspirin and 46.7 parts of sodium bicarbonate, the amount of water is about 50 parts. Therefore, the amount of aspirin is about 50% of the total, the amount of the bicarbonate is about 25% and the amount of water is 25% also. The purpose of using such high concentrations is to utilize the equipment capacity to its maximum, to produce the maximum yield on crystallization and to use the smallest amount of the solvent. Also, the use of high concentrations causes the reaction to be completed in shorter time, since the rate of a chemical reaction is proportional to the product of the concentrations of the reactants. In the present case, the rate can be expressed by the equation: EQU R=C.sub.1 .times.C.sub.2
where R is the rate and C.sub.1 and C.sub.2 are the concentrations of aspirin and sodium bicarbonate, respectively.
The concentration of both reactants varies constantly as the reaction proceeds. The initial concentration of aspirin is low because of its low solubility, whereas that of the more soluble sodium bicarbonate is about 33%.
The situation is quite different when the use of single doses by individual patients is considered. Aspirin tablets are taken with a half-glass of water about 100 to 120 ml; about 31/2 to 4 oz). While the concentration of aspirin is the same in any amount of water, its value being determined by its solubility in water (0.33%) and is thus constant, the concentration of sodium bicarbonate (C.sub.2 in the above equation) can be varied as desired within relatively wide limits. However, if one wishes to use it in equimolecular proportions, a 325 mg dose of aspirin will require 152 mg of sodium bicarbonate to produce a solution. If this dose is taken in 100 ml of water, the concentration of sodium bicarbonate will be, initially, 0.152% and will decrease as the reaction progresses. Thus, concentration values in the laboratory or the plant, on one hand, and in personal usage on the other, are 33% vs. 0.15%. Referring to the equation above, it is evident that the rate of the reaction in the later case will be very much lower than in the former.
In order to bring the rate of the reaction within practical limits, and since it is not possible to increase the concentration of aspirin, the only alternative is to increase the concentration of sodium well beyond the stoichiometric proportions. As pointed out, the sodium content of such soluble aspirin products is so high as to make them unsuitable for most of the major applications of aspirin in medicine.
The influence of sodium bicarbonate concentrations on its rate of reaction with aspirin was determined by stirring 325 mg of aspirin with variable amounts of sodium bicarbonate in 100 ml of water and recording the time needed for the formation of a solution. Aspirin USP mesh #325 was used. This is the finest particle size available commercially (Monsanto, "micronized"). In order to simulate the conditions of practical use as closely as possible, the mixture gas stirred by hand and with a teaspoon in an 8 oz glass.
TABLE 1 ______________________________________ sodium bicarbonate (mg) 2,000 1,000 500 250 152 reaction time (minutes) 0.5 1.0 2.2 5.5 12 ______________________________________
In order to be of practical use to individual patients, a dose of a soluble aspirin product should dissolve in about half a glass of water (100-120 ml; 3- 1/2-4 oz), in a reasonably short time (less than 60 seconds), with stirring by hand with a spoon. As mentioned, this is achieved in commercial products by the use of a large excess of sodium bicarbonate. The necessity for such an excess is also seen when Table 1 is examined. However, the price is the introduction of amounts of sodium that are unacceptable, if the product is to be used on a regular basis. As can also be seen from Table 1, if the sodium bicarbonate is reduced to its theoretical amount (the minimum necessary for the reaction is 152 mg) the reaction rate becomes unacceptably low (12 minutes).